History of the periodic table
Part of a series on the |
Periodic table |
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The
The history of the periodic table reflects over two centuries of growth in the understanding of the chemical and physical properties of the elements, with major contributions made by
Early history
Nine chemical elements – carbon, sulfur, iron, copper, silver, tin, gold, mercury, and lead, have been known since before antiquity, as they are found in their native form and are relatively simple to mine with primitive tools.[3] Around 330 BCE, the Greek philosopher Aristotle proposed that everything is made up of a mixture of one or more roots, an idea originally suggested by the Sicilian philosopher Empedocles. The four roots, which the Athenian philosopher Plato called elements, were earth, water, air and fire. Similar ideas about these four elements existed in other ancient traditions, such as Indian philosophy.
A few extra elements were known in the age of alchemy:
First categorizations
The history of the periodic table is also a history of the
In 1661, Boyle defined elements as "those primitive and simple Bodies of which the mixt ones are said to be composed, and into which they are ultimately resolved."[6]
In 1718, Étienne François Geoffroy's Affinity Table made use of several aspects — (1) tabular grouping and (2) correlation with chemical affinity — that would later be reprised.
In 1789, French chemist Antoine Lavoisier wrote Traité Élémentaire de Chimie (Elementary Treatise of Chemistry), which is considered to be the first modern textbook about chemistry. Lavoisier defined an element as a substance whose smallest units cannot be broken down into a simpler substance.[8] Lavoisier's book contained a list of "simple substances" that Lavoisier believed could not be broken down further, which included oxygen, nitrogen, hydrogen, phosphorus, mercury, zinc and sulfur, which formed the basis for the modern list of elements. Lavoisier's list also included "light" and "caloric", which at the time were believed to be material substances. He classified these substances into metals and nonmetals. While many leading chemists refused to believe Lavoisier's new revelations, the Elementary Treatise was written well enough to convince the younger generation. However, Lavoisier's descriptions of his elements lack completeness, as he only classified them as metals and non-metals.
In 1808–10, British natural philosopher
In 1815, British physician and chemist William Prout noticed that atomic weights seemed to be multiples of that of hydrogen.[9][10]
In 1817, German physicist Johann Wolfgang Döbereiner began to formulate one of the earliest attempts to classify the elements.[11] In 1829, he found that he could form some of the elements into groups of three, with the members of each group having related properties. He termed these groups triads.[12]
Definition of Triad law
"Chemically analogous elements arranged in increasing order of their atomic weights formed well marked groups of three called Triads in which the atomic weight of the middle element was found to be generally the arithmetic mean of the atomic weight of the other two elements in the triad.
- chlorine, bromine, and iodine
- calcium, strontium, and barium
- sulfur, selenium, and tellurium
- lithium, sodium, and potassium"
All those attempts to sort elements by atomic weights were inhibited by the inaccurate determination of weights, and not just slightly: carbon, oxygen and many other elements were believed to be half their actual masses (cf. the illustration by Dalton above), because only monatomic gases were believed to exist.[13] Even though Amedeo Avogadro and, independently of him, André-Marie Ampère, proposed the solution in the form of diatomic molecules and Avogadro's law already in the 1810s, it was not until after Stanislao Cannizzaro's publications in late 1850s when the theory began to be widely considered.
In 1860, the modern scientific consensus emerged at the first international chemical conference, the Karlsruhe Congress, and a revised list of elements and atomic masses was adopted. It helped spur creation of more extensive systems. The first such system emerged in two years.[14]
Comprehensive formalizations
Properties of the elements, and thus properties of light and heavy bodies formed by them, are in a periodic dependence on their atomic weight.
— Russian chemist Dmitri Mendeleev, formulating the periodic law for the first time in his 1871 article "Periodic regularity of the chemical elements"[15]
French geologist Alexandre-Émile Béguyer de Chancourtois noticed that the elements, when ordered by their atomic weights, displayed similar properties at regular intervals. In 1862, he devised a three-dimensional chart, named the "telluric helix", after the element tellurium, which fell near the center of his diagram.[16][17] With the elements arranged in a spiral on a cylinder by order of increasing atomic weight, de Chancourtois saw that elements with similar properties lined up vertically. The original paper from Chancourtois in Comptes rendus de l'Académie des Sciences did not include a chart and used geological rather than chemical terms. In 1863, he extended his work by including a chart and adding ions and compounds.[18]
The next attempt was made in 1864. British chemist John Newlands presented in Chemical News[19] a classification of the 62 known elements. Newlands noticed recurring trends in physical properties of the elements at recurring intervals of multiples of eight in order of mass number;[20] based on this observation, he produced a classification of these elements into eight groups. Each group displayed a similar progression; Newlands likened these progressions to the progression of notes within a musical scale.[17][21][22][23] Newlands's table left no gaps for possible future elements, and in some cases had two elements at the same position in the same octave. Newlands's table was ignored or ridiculed by some of his contemporaries.[19] The Chemical Society refused to publish his work. The president of the Society, William Odling, defended the Society's decision by saying that such "theoretical" topics might be controversial;[24] there was even harsher opposition from within the Society, suggesting the elements could have been just as well listed alphabetically.[14] Later that year, Odling suggested a table of his own[25] but failed to get recognition following his role in opposing Newlands's table.[24]
German chemist
In 1869, Russian chemist
-
Meyer's periodic table, published in "Die modernen Theorien der Chemie", 1864[26]
-
Newlands's law of octaves, 1866
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Mendeleev's first Attempt at a system of elements, 1869
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Mendeleev's Natural system of the elements, 1870
-
Mendeleev's periodic table, 1871
Priority dispute and recognition
That person is rightly regarded as the creator of a particular scientific idea who perceives not merely its philosophical, but its real aspect, and who understands so to illustrate the matter so that everyone can become convinced of its truth. Then alone the idea, like matter, becomes indestructible.
— Mendeleev in his 1881 article in British journal Chemical News in a correspondence debate with Meyer over priority of the periodic table invention[34]
Mendeleev's predictions and inability to incorporate the rare-earth metals
Name | Atomic weight | Modern name (year of discovery) | |
---|---|---|---|
Mendeleev | Modern | ||
Ether | 0.17 | — | — |
Coronium | 0.4 | — | — |
Eka-boron | 44 | 44.6 | Scandium |
Eka-cerium | 54 | — | — |
Eka-aluminum | 68 | 69.2 | Gallium |
Eka-silicon | 72 | 72.0 | Germanium |
Eka-manganese | 100 | 99 | Technetium (1925) |
Eka-molybdenum | 140 | — | — |
Eka-niobium | 146 | — | — |
Eka-cadmium | 155 | — | — |
Eka-iodine | 170 | — | — |
Tri-manganese | 190 | 186 | Rhenium (1925) |
Eka-caesium | 175 | — | — |
Dvi-tellurium | 212 | 210 | Polonium (1898) |
Dvi-caesium | 220 | 223 | Francium (1937) |
Eka-tantalum | 235 | 231 | Protactinium (1917) |
Even as Mendeleev corrected positions of some elements, he thought that some relationships that he could find in his grand scheme of periodicity could not be found because some elements were still undiscovered, and that the properties of such undiscovered elements could be deduced from their expected relationships with other elements. In 1870, he first tried to characterize the yet undiscovered elements, and he gave detailed predictions for three elements, which he termed eka-boron, eka-aluminium, and eka-silicium;[37] he also more briefly noted a few other expectations.[38] It has been proposed that the prefixes eka, dvi, and tri, Sanskrit for one, two, and three, respectively, are a tribute to Pāṇini and other ancient Sanskrit grammarians for their invention of a periodic alphabet.[28] In 1871, Mendeleev expanded his predictions further.
Compared to the rest of the work, Mendeleev's 1869 list misplaces seven then known elements: indium, thorium, and five rare-earth metals: yttrium, cerium, lanthanum, erbium, and didymium. The last two were later found to be mixtures of two different elements; ignoring those would allow him to restore the logic of increasing atomic weight. These elements (all thought to be divalent at the time) puzzled Mendeleev as they did not show a regular increase in valency despite their seemingly consequential atomic weights.[39] Mendeleev grouped them together, thinking of them as of a particular kind of series.[c] In early 1870, he decided that the weights for these elements must be wrong and that the rare-earth metals should be trivalent (which accordingly increased their predicted atomic weights by half). He measured the heat capacity of indium, uranium, and cerium to demonstrate their higher assumed valency (which was soon confirmed by Prussian chemist Robert Bunsen).[40] Mendeleev treated the change by assessing each element to an individual place in his system of the elements rather than continuing to treat them as a series.
Mendeleev noticed that there was a significant difference in atomic mass between cerium and
In addition to the predictions of scandium, gallium, and germanium that were quickly realized, Mendeleev's 1871 table left many more spaces for undiscovered elements, though he did not provide detailed predictions of their properties. In total, he predicted eighteen elements, though only half corresponded to elements that were later discovered.[45]
Priority of discovery
None of the proposals were accepted immediately, and many contemporary chemists found it too abstract to have any meaningful value. Of those chemists that proposed their categorizations, Mendeleev strove to back his work and promote his vision of periodicity, Meyer did not promote his work very actively, and Newlands did not make a single attempt to gain recognition abroad.[citation needed]
Both Mendeleev and Meyer created their respective tables for their pedagogical needs; the difference between their tables is well explained by the fact that the two chemists sought to use a formalized system to solve different problems.[46] Mendeleev's intent was to aid composition of his textbook, Foundations of Chemistry, whereas Meyer was rather concerned with presentation of theories.[46] Mendeleev's predictions emerged outside of the pedagogical scope in the realm of journal science,[47] while Meyer made no predictions at all and explicitly stated his table and his textbook it was contained in, Modern Theories, should not be used for prediction in order to make the point to his students to not make too many purely theoretically constructed projections.[48]
Mendeleev and Meyer differed in temperament, at least when it came to promotion of their respective works. Boldness of Mendeleev's predictions was noted by some contemporary chemists, however skeptical they may have been.[49] Meyer referred to Mendeleev's "boldness" in an edition of Modern Theories, whereas Mendeleev mocked Meyer's indecisiveness to predict in an edition of Foundations of Chemistry.[49]
Recognition of Mendeleev's table
Eventually, the periodic table was appreciated for its descriptive power and for finally systematizing the relationship between the elements,[50] although such appreciation was not universal.[51] In 1881, Mendeleev and Meyer had an argument via an exchange of articles in British journal Chemical News over priority of the periodic table, which included an article from Mendeleev, one from Meyer, one of critique of the notion of periodicity, and many more.[52] In 1882, the Royal Society in London awarded the Davy Medal to both Mendeleev and Meyer for their work to classify the elements; although two of Mendeleev's predicted elements had been discovered by then, Mendeleev's predictions were not at all mentioned in the prize rationale.
Mendeleev's eka-aluminium was discovered in 1875 and became known as gallium; eka-boron and eka-silicium were discovered in 1879 and 1886, respectively, and were named scandium and germanium.[17] Mendeleev was even able to correct some initial measurements with his predictions, including the first prediction of gallium, which matched eka-aluminium fairly closely but had a different density. Mendeleev advised the discoverer, French chemist Paul-Émile Lecoq de Boisbaudran, to measure the density again; de Boisbaudran was initially skeptical (not least because he thought Mendeleev was trying to take credit from him) but eventually admitted the correctness of the prediction. Mendeleev contacted all three discoverers; all three noted the close similarity of their discovered elements with Mendeleev's predictions, with the last of them, German chemist Clemens Winkler, admitting this suggestion was not first made by Mendeleev or himself after the correspondence with him, but by a different person, German chemist Hieronymous Theodor Richter.[e] Some contemporary chemists were not convinced by these discoveries, noting the dissimilarities between the new elements and the predictions or claiming those similarities that did exist were coincidental.[51] However, success of Mendeleev's predictions helped spread the word about his periodic table.[54] Later, chemists used the successes of these Mendeleev's predictions to justify his table.[14]
By 1890, Mendeleev's periodic table had been universally recognized as a piece of basic chemical knowledge.[55] Apart from his own correct predictions, a number of aspects may have contributed to this, such as the correct accommodation of many elements whose atomic weights were thought to have wrong values but were later corrected.[54] The debate on the position of the rare-earth metals helped spur the discussion about the table as well.[54][f] In 1889, Mendeleev noted at the Faraday Lecture to the Royal Institution in London that he had not expected to live long enough "to mention their discovery to the Chemical Society of Great Britain as a confirmation of the exactitude and generality of the periodic law".[56]
Inert gases and ether
The great value of Newland's, Mendeleef's, and Lothar Meyer's generalisation, known as the periodic arrangement of the elements, is universally acknowledged. But a study of this arrangement, it must be allowed, is a somewhat tantalising pleasure; for, although the properties of elements do undoubtedly vary qualitatively, and, indeed, show approximate quantitative relations to their position in the periodic table, yet there are inexplicable deviations from regularity, which hold forth hopes of the discovery of a still more far-reaching generalisation. What that generalisation may be is not yet to be divined; but that it must underlie what is known, and must furnish a clue to the explanation of irregularities, cannot be disputed.
— British chemists William Ramsay and Morris Travers in 1900 discussion of their research of new inert gases[57]
Inert gases
British chemist Henry Cavendish, the discoverer of hydrogen in 1766, discovered that air is composed of more gases than nitrogen and oxygen.[58] He recorded these findings in 1784 and 1785; among them, he found a then-unidentified gas less reactive than nitrogen. Helium was first reported in 1868; the report was based on the new technique of spectroscopy; some spectral lines in light emitted by the Sun did not match those of any of the known elements. Mendeleev was not convinced by this finding since variance of temperature led to change of intensity of spectral lines and their location on the spectrum;[59] this opinion was held by some other scientists of the day. Others believed the spectral lines could belong to an element that occurred on the Sun but not on Earth; some believed it was yet to be found on Earth.[citation needed]
In 1894, British chemist
In 1896, Ramsay tested a report of American chemist
Changes to the periodic table
Although the sequence of atomic weights suggested that inert gases should be located between halogens and alkali metals, and there were suggestions to put them into group VIII coming from as early as 1895,[70] such placement contradicted one of Mendeleev's basic considerations, that of the highest oxides. Inert gases did not form any oxides, and no other compounds at all, and as such, their placement in a group where elements should form tetroxides was seen as merely auxiliary and not natural; Mendeleev doubted inclusion of those elements in group VIII.[70] Later developments, particularly by British scientists, focused on correspondence of inert gases with halogens to their left and alkali metals to their right. In 1898, when only helium, argon, and krypton were definitively known, Crookes suggested these elements be placed in a single column between the hydrogen group and the fluorine group.[71] In 1900, at the Prussian Academy of Sciences, Ramsay and Mendeleev discussed the new inert gases and their location in the periodic table; Ramsay proposed that these elements be put in a new group between halogens and alkali metals, to which Mendeleev agreed.[54] Ramsay published an article after his discussions with Mendeleev; the tables in it featured halogens to the left of inert gases and alkali metals to the right.[72] Two weeks before that discussion, Belgian botanist Léo Errera had proposed to the Royal Academy of Science, Letters and Fine Arts of Belgium to put those elements in a new group 0. In 1902, Mendeleev wrote that those elements should be put in a new group 0; he said this idea was consistent with what Ramsay suggested to him and referred to Errera as to the first person to suggest the idea.[73] Mendeleev himself added these elements to the table as group 0 in 1902, without disturbing the basic concept of the periodic table.[73][74]
In 1905, Swiss chemist Alfred Werner resolved the dead zone of Mendeleev's table. He determined that the rare-earth elements (lanthanides), 13 of which were known, lay within that gap. Although Mendeleev knew of lanthanum, cerium, and erbium, they were previously unaccounted for in the table because their total number and exact order were not known; Mendeleev still could not fit them in his table by 1901.[68] This was in part a consequence of their similar chemistry and the imprecise determination of their atomic masses. Combined with the lack of a known group of similar elements, this rendered the placement of the lanthanides in the periodic table difficult.[75] This discovery led to a restructuring of the table and the first appearance of the 32-column form.[76]
Ether
By 1904, Mendeleev's table rearranged several elements, and included the noble gases along with most other newly discovered elements. It still had the dead zone, and a row zero was added above hydrogen and helium to include
Mendeleev was not satisfied with the lack of understanding of the nature of this periodicity; this would only be possible with the understanding of composition of atom. However, Mendeleev firmly believed that future would only develop the notion rather than challenge it and reaffirmed his belief in writing in 1902.[78]
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Main table of the periodic table published by Australian chemist David Orme Masson in 1895
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Fragment of a periodic table published by Ramsay in 1896
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Fragment of a periodic table published by Ramsay in 1900
-
Periodic table as published by Errera in 1900
-
Werner's 32-column 1905 table. This table left spaces for many then-unknown elements, and several elements had their positions revised following advances in atomic theory.
Atomic theory and isotopes
Radioactivity and isotopes
In 1907 it was discovered that thorium and radiothorium, products of radioactive decay, were physically different but chemically identical; this led Frederick Soddy to propose in 1910 that they were the same element but with different atomic weights.[79] Soddy later proposed to call these elements with complete chemical identity "isotopes".[80]
The problem of placing isotopes in the periodic table had arisen beginning in 1900 when four radioactive elements were known: radium, actinium, thorium, and uranium. These radioactive elements (termed "radioelements") were accordingly placed at the bottom of the periodic table, as they were known to have greater atomic weights than stable elements, although their exact order was not known. Researchers believed there were still more radioactive elements yet to be discovered, and during the next decade, the decay chains of thorium and uranium were extensively studied. Many new radioactive substances were found, including the noble gas radon, and their chemical properties were investigated.[17] By 1912, almost 50 different radioactive substances had been found in the decay chains of thorium and uranium. American chemist Bertram Boltwood proposed several decay chains linking these radioelements between uranium and lead. These were thought at the time to be new chemical elements, substantially increasing the number of known "elements" and leading to speculations that their discoveries would undermine the concept of the periodic table which had long been established to obey the octet rule.[45] For example, there was not enough room between lead and uranium to accommodate these discoveries, even assuming that some discoveries were duplicates or incorrect identifications. It was also believed that radioactive decay violated one of the central principles of the periodic table, namely that chemical elements could not undergo transmutations and always had unique identities.[17]
Soddy and Kazimierz Fajans, who had been following these developments, published in 1913 that although these substances emitted different radiation,[81] many of these substances were identical in their chemical characteristics, so shared the same place in the periodic table.[82][83] They became known as isotopes, from the Greek isos topos ("same place").[17][84] Austrian chemist Friedrich Paneth cited a difference between "real elements" (elements) and "simple substances" (isotopes), also determining that the existence of different isotopes was mostly irrelevant in determining chemical properties.[45]
Following British physicist
Rutherford model and atomic number
In 1913, amateur Dutch physicist Antonius van den Broek was the first to propose that the atomic number (nuclear charge) determined the placement of elements in the periodic table. He correctly determined the atomic number of all elements up to atomic number 50 (tin), though he made several errors with heavier elements. However, Van den Broek did not have any method to experimentally verify the atomic numbers of elements; thus, they were still believed to be a consequence of atomic weight, which remained in use in ordering elements.[85]
Moseley was determined to test Van den Broek's hypothesis.
Swedish physicist
Electron shell and quantum mechanics
In 1888,
The chemist Charles Rugeley Bury made the next major step toward our modern theory in 1921, by suggesting that eight and eighteen electrons in a shell form stable configurations. Bury's scheme was built upon that of earlier chemists and was a chemical model. Bury proposed that the electron configurations in transitional elements depended upon the valency electrons in their outer shell.[94] In some early papers, the model was called the "Bohr-Bury Atom". He introduced the word transition to describe the elements now known as transition metals or transition elements.[95]
In the 1910s and 1920s, pioneering research into quantum mechanics led to new developments in atomic theory and small changes to the periodic table. In the 19th century, Mendeleev had already asserted that there was a fixed periodicity of eight, and expected a mathematical correlation between atomic number and chemical properties.[96] The Bohr model was developed beginning 1913, and championed the idea of electron configurations that determine chemical properties. Bohr proposed that elements in the same group behaved similarly because they have similar electron configurations, and that noble gases had filled valence shells;[97] this forms the basis of the modern octet rule. Bohr's study of spectroscopy and chemistry was not usual among theoretical atomic physicists. Even Rutherford told Bohr that he was struggling "to form an idea of how you arrive at your conclusions".[98] This is because none of the quantum mechanical equations describe the number of electrons per shell and orbital. Bohr acknowledged that he was influenced by the work of Walther Kossel, who in 1916 was the first to establish an important connection between the quantum atom and the periodic table. He noticed that the difference between the atomic numbers 2, 10, 18 of the first three noble gases, helium, neon, argon, was 8, and argued that the electrons in such atoms orbited in "closed shells". The first contained only 2 electrons, the second and third, 8 each.[99][100] Bohr's research then led Austrian physicist Wolfgang Pauli to investigate the length of periods in the periodic table in 1924. Pauli demonstrated that this was not the case. Instead, the Pauli exclusion principle was developed, not upon a mathematical basis, but upon the previous developments in alignment with chemistry.[101] This rule states that no electrons can coexist in the same quantum state, and showed, in conjunction with empirical observations, the existence of four quantum numbers and the consequence on the order of shell filling.[97] This determines the order in which electron shells are filled and explains the periodicity of the periodic table.
British chemist Charles Bury is credited with the first use of the term transition metal in 1921 to refer to elements between the main-group elements of groups II and III. He explained the chemical properties of transition elements as a consequence of the filling of an inner subshell rather than the valence shell. This proposition, based upon the work of American chemist Gilbert N. Lewis, suggested the appearance of the d subshell in period 4 and the f subshell in period 6, lengthening the periods from 8 to 18 and then 18 to 32 elements, thus explaining the position of the lanthanides in the periodic table.[102]
Proton and neutron
The discovery of proton and neutron demonstrated that an atom was divisible; this rendered Lavoisier's definition of a chemical element obsolete. A chemical element is defined today as a species of atoms with a consistent number of protons[103] and that number is now known to be precisely the atomic number of an element. The discovery also explained the mechanism of several types of radioactive decay, such as alpha decay.
Eventually, it was proposed that protons and neutrons were made of even smaller particles called quarks; their discovery explained the transmutation of neutrons into protons in beta decay.
From short form into long form (into -A and -B groups)
Circa 1925, the periodic table changed by shifting some Reihen (series) to the right, into an extra set of columns (groups). The original groups I–VII were repeated, distinguished by adding "A" and "B". Group VIII (with three columns) remained sole.
Thus, Reihen 4 and 5 were shifted, and together formed new period 4 with groups IA–VIIA, VIII, IB–VIIB.
modern (long): | IUPAC group | 1 | 2 | no number | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 | × | 11 | 12 | 13 | 14 | 15 | 16 | 17 | 18 | |
1900+ (long): | old IUPAC (A–B, Europe) | IA | IIA | IIIA | IVA | VA | VIA | VIIA | VIII | IB | IIB | IIIB | IVB | VB | VIB | VIIB | 0 | |||||
CAS (A–B–A, US) | IA | IIA | IIIB | IVB | VB | VIB | VIIB | VIIIB | IB | IIB | IIIA | IVA | VA | VIA | VIIA | VIIIA | ||||||
1871 (short)→ | Gruppe | I | II | [ ] | III | IV | V | VI | VII | VIII | I |
II |
III |
IV |
V |
VI |
VII |
0 | ||||
Period ① | Reihe 1 | H | He | |||||||||||||||||||
Period ② | Reihe 2 | Li | Be | B | C | N | O | F | Ne | |||||||||||||
Period ③ | Reihe 3 | Na | Mg | Al | Si | P | S | Cl | Ar | |||||||||||||
Period ④ | Reihe 4 | K | Ca | –Sc | Ti | V | Cr | Mn | Fe | Co | Ni | Cu(1st) | ||||||||||
Reihe 5 | (Cu)(2nd) | Zn | – Ga | – Ge | As | Se | Br | Kr | ||||||||||||||
Period ⑤ | Reihe 6 | Rb | Sr | Yt[=Y] | Zr | Nb | Mo | –([Tc]) | Ru | Rh | Pd | Ag(1st) | ||||||||||
Reihe 7 | (Ag)(2nd) | Cd | In | Sn | Sb | Te | J[=I] | Xe | ||||||||||||||
Period ⑥ | Reihe 8 | Cs | Ba | [La–Lu] | — | — | — | — | — | — | —(1st) | |||||||||||
Reihe 9 | (—)(2nd) | — | — | — | — | — | — | |||||||||||||||
Period ⑥ | Reihe 10 | — | — | Ta | W | –([Re]) | Os | Ir | Pt | Au(1st) | ||||||||||||
Reihe 11 | (Au)(2nd) | Hg | Tl | Pb | Bi | —[Po] | —[At] | [Rn] | ||||||||||||||
Period ⑦ | Reihe 12 | —[Fr] | —Rd[=Ra] | [Ac–Lr] | — |
— |
— | — | — | — | — | |||||||||||
Bold text | in Periodic Table 1871 |
Italic text | in Periodic Table 1906 (the last by Mendeleev) |
Regular text (not bold) | added after 1906 |
begin–end of 1871 Reihe | |
– Ga | Element predicted, later proven correct within Mendeleev's lifetime and added by him |
– (Tc) | Element predicted, later proven correct posthumously |
– | Element projected, but not predicted |
Element predicted, later proven wrong due to – not an element ("?Di"), or –wrong position (" | |
[ ] | Added or changed after 1871 |
Cu(1st) × / (Cu)(2nd) | Element mentioned twice: in Gruppe VIII and I. The 2nd mentioning survived, Gruppe/group VIII was reduced from four columns to three (×) |
Published 1871, English version: "Reihen" translated as "Series" (that is, arrays with regularity not just rows). Reproduced in Scerri (2007), p. 111 |
Later expansions and the end of the periodic table
We already feel that we have neared the moment when this [periodic] law begins to change, and change fast.
Actinides
As early as 1913, Bohr's research on
In 1940,
Seaborg's
In light of these observations and an apparent explanation for the chemistry of transuranic elements, and despite fear among his colleagues that it was a radical idea that would ruin his reputation, Seaborg nevertheless submitted it to Chemical & Engineering News and it gained widespread acceptance; new periodic tables thus placed the actinides below the lanthanides.[107] Following its acceptance, the actinide concept proved pivotal in the groundwork for discoveries of heavier elements, such as berkelium in 1949.[109] It also supported experimental results for a trend towards +3 oxidation states in the elements beyond americium—a trend observed in the analogous 4f series.[105]
Relativistic effects and expansions beyond period 7
Seaborg's subsequent elaborations of the actinide concept theorized a series of
The discovery of tennessine in 2010 filled the last remaining gap in the seventh period. Any newly discovered elements will thus be placed in an eighth period.
Despite the completion of the seventh period, experimental chemistry of some transactinides has been shown to be inconsistent with the periodic law. In the 1990s, Ken Czerwinski at University of California, Berkeley observed similarities between rutherfordium and plutonium and between dubnium and protactinium, rather than a clear continuation of periodicity in groups 4 and 5. More recent experiments on copernicium and flerovium have yielded inconsistent results, some of which suggest that these elements behave more like the noble gas radon rather than mercury and lead, their respective congeners. As such, the chemistry of many superheavy elements has yet to be well characterized, and it remains unclear whether the periodic law can still be used to extrapolate the properties of undiscovered elements.[2][113]
See also
- History of chemistry
- Periodic systems of small molecules
- The Mystery of Matter: Search for the Elements (PBS film)
- Discovery of chemical elements
- Types of periodic tables
Notes
- ^ They were R2O, R2O2, R2O3, R2O4, R2O5, R2O6, and R2O7. The list was later appended with R2O8.
- ^ Scerri notes that this table "does not include elements such as astatine and actinium, which he [Mendeleev] predicted successfully but did not name. Neither does it include predictions that were represented just by dashes in Mendeleev’s periodic systems. Among some other failures, not included in the table, is an inert gas element between barium and tantalum, which would have been called ekaxenon, although Mendeleev did not refer to it as such."[36]
- ^ He noted similarity despite sequential atomic weights; he termed such sequences as primary groups (as opposed to regular secondary groups, such as the halogens and the alkali metals). Other primary groups were rhodium, ruthenium, and palladium; and iridium, osmium, and platinum.
- ^ Mendeleev referred to Brauner in this manner after Brauner measured the atomic weight of tellurium and obtained the value 125. Mendeleev had thought that due to the properties tellurium and iodine display, the latter should be the heavier one while the contemporary data pointed otherwise (tellurium was assessed with the value of 128, and iodine 127). Later measurements by Brauner himself, however, showed the correctness of the original measurement; Mendeleev doubted it for the rest of his life.[44]
- ^ Notably, Mendeleev did not immediately identify germanium as eka-silicium. Winkler explained, "The present case, however, shows quite clearly how deceptive it can be to use analogies, because the tetradic value of germanium has meanwhile become an irrefutable fact, and there can be no doubt that the new element is nothing other than "eka-silicium" predicted by Mendeleev fifteen years ago. This identification comes from the short and still very imperfect characteristic of germanium that I gave at the beginning and was first decisively pronounced by V. v. Richter. Almost at the same time, Mendeleev, the deserving creator of the periodic system, commented that although several of the properties of germanium I mentioned reminded of those of eka-silicium, the observed liquidity of the element indicated the possibility of placing it elsewhere in the periodic system. Lothar Meyer declared the germanium to be eka-silicium from the beginning, adding that according to the atomic volume curve produced by it, contrary to Mendeleev's assumption, it had to be easily meltable and probably also easy to vaporize. At that time the germanium had not yet been presented in the reguline state; it is all the more remarkable that, as will be shown below, Lothar Meyer's condition has, to some extent, really come true."[53]
- ^ Meyer's tables, in contrast, did not at all attempt to incorporate those elements.[citation needed]
- ^ The only other monatomic gas known at the time was vaporized mercury.[60]
- ^ Mendeleev did consider that some atomic weight values could be missing from the set of known values. However, Mendeleev could not have made a prediction of a group of unreactive gases in a fashion similar to the one in which he made his predictions on reactive elements and their chemical properties.[69]
- special theory of relativity; the idea that ether did not exist was accepted in the scientific community rather quickly.
References
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- ^ Prout, William (November 1815). "On the relation between the specific gravities of bodies in their gaseous state and the weights of their atoms". Annals of Philosophy. 6: 321–330.
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- . Here, Döbereiner found that strontium's properties were intermediate to those of calcium and barium.
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- ^ John Newlands, Chemistry Review, November 2003, pp. 15-16.[full citation needed]
- ^ See:
- Newlands, John A. R. (7 February 1863). "On relations among the equivalents". The Chemical News. 7: 70–72.
- Newlands, John A. R. (30 July 1864). "Relations between equivalents". The Chemical News. 10: 59–60.
- Newlands, John A. R. (20 August 1864). "On relations among the equivalents". The Chemical News. 10: 94–95.
- Newlands, John A. R. (18 August 1865). "On the law of octaves". The Chemical News. 12: 83.
- (Editorial staff) (9 March 1866). "Proceedings of Societies: Chemical Society: Thursday, March 1". The Chemical News. 13: 113–114.
- Newlands, John A.R. (1884). On the Discovery of the Periodic Law and on Relations among the Atomic Weights. E. & F.N. Spon: London, England.
- ^ in a letter published in Chemistry News in February 1863, according to the Notable Names Data Base
- ^ "An Unsystematic Foreshadowing: J. A. R. Newlands". web.lemoyne.edu. Retrieved 2019-07-13.
- ^ ISBN 9783642283857. From p. 38: "The reason [for rejecting Newlands's paper, which was] given by Odling, then the president of the Chemical Society, was that they made a rule not to publish theoretical papers, and this on the quite astonishing grounds that such papers lead to a correspondence of controversial character."
- ^ See:
- Odling, William (June 1857). "On the natural groupings of the elements. Part 1". Philosophical Magazine. 4th series. 13 (88): 423–440. .
- Odling, William (1857). "On the natural groupings of the elements. Part 2". Philosophical Magazine. 4th series. 13 (89): 480–497. .
- Odling, William (1864). "On the hexatomicity of ferricum and aluminium". Philosophical Magazine. 4th series. 27 (180): 115–119. .
- Odling, William (1864). "On the proportional numbers of the elements". Quarterly Journal of Science. 1: 642–648.
- ^ a b Meyer, Julius Lothar; Die modernen Theorien der Chemie (1864); table on page 137.
- ISBN 0-03-073168-2
- ^ S2CID 209975833.
- ^ Mendeleev, Dmitri (1869). "Versuche eines Systems der Elemente nach ihren Atomgewichten und chemischen Functionen" [System of Elements according to their Atomic Weights and Chemical Functions]. Journal für Praktische Chemie. 106: 251.
- ^ Менделеев, Д. (1869). "Соотношение свойств с атомным весом элементов" [Relationship of properties of the elements to their atomic weights]. Журнал Русского Химического Общества (Journal of the Russian Chemical Society) (in Russian). 1: 60–77.
- ^ Mendeleev, Dmitri (1869). "Ueber die Beziehungen der Eigenschaften zu den Atomgewichten der Elemente" [On the relations of properties of the elements to their atomic weights]. Zeitschrift für Chemie. 12: 405–406.
- ^ Petrov 1981, p. 65.
- ^ Mendeleev 1870, p. 76.
- ^ Scerri 2019, p. 147.
- ^ Scerri 2019, p. 142.
- ^ Scerri 2019, p. 143.
- ^ Mendeleev 1870, pp. 90–98.
- ^ Mendeleev 1870, pp. 98–101.
- ^ Thyssen & Binnemans 2015, p. 159.
- ^ Thyssen & Binnemans 2015, pp. 174–175.
- S2CID 59564667.
- ^ Thyssen & Binnemans 2015, p. 177.
- ^ Thyssen & Binnemans 2015, pp. 179–181.
- ^ Scerri 2019, pp. 130–131.
- ^ .
- ^ a b Gordin 2012, pp. 75–76.
- ^ Gordin 2012, p. 76.
- ^ Gordin 2012, pp. 71–74.
- ^ a b Gordin 2012, p. 75.
- JSTOR 26057945.
- ^ a b Scerri 2019, pp. 170–172.
- ^ Scerri 2019, pp. 147–149.
- .
- ^ a b c d Scerri 2019, p. 156.
- ^ Scerri 2019, p. 157.
- ^ Rouvray, R. "Dmitri Mendeleev". New Scientist. Retrieved 2020-04-19.
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- ^ Assovskaya, A. S. (1984). "Первый век гелия" [The first century of helium]. Гелий на Земле и во Вселенной [Helium on Earth and in the Universe] (in Russian). Leningrad: Nedra.
- ^ Scerri 2019, p. 151.
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- ^ Petrov 1981, pp. 38–44.
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- ^ Petrov 1981, pp. 59–61.
- ^ Petrov 1981, pp. 54–55.
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{{cite book}}
: CS1 maint: location missing publisher (link)
An English translation appeared as
Mendeléeff, D. (1904). An Attempt Towards A Chemical Conception Of The Ether. Translated by G. Kamensky. Longmans, Green & Co. - ISBN 978-0-470-01005-1.
- ^ S2CID 104132201.
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- ^ Trifonov, D. N. "Д.И. Менделеев. Нетрадиционный взгляд (II)" [D.I. Mendeleev. An unconventional view (II)] (in Russian). Moscow State University. Retrieved 2020-04-12.
- ^ Howorth, Muriel (1958), The Life of Frederick Soddy (London: New World)
- ^ Soddy, Frederick (1913), 'Intra-Atomic Charge', Nature, 92, 399–400
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- .
- ^ Soddy, Frederick (28 February 1913). "The radio-elements and the periodic law". The Chemical News. 107 (2779): 97–99.
- S2CID 3965303. See p. 400.
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- ^ J.R. Rydberg, Den Kungliga Svenska Vetenskapsakadem- iens Handlingar 23 (11) (1889).
- Octet Rule
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- .
- ^ I. Langmuir, "The Arrangement of Electrons in Atoms and Molecules", J. Am. Chem. Soc., 41 (1919), p.868-934.
- ^ C.R. Bury, “Langmuir’s Theory on the Arrangement of Electrons in Atoms and Molecules, J. Am. Chem. Soc., 43 (1921), 1602-1609.
- .
The first use of the term "transition" in its modern electronic sense appears to be due to the British chemist C. R.Bury, who first used the term in his 1921 paper on the electronic structure of atoms and the periodic table
- S2CID 118775532.
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- ^ Pais, Abraham (1991), Niels Bohr’s Times, in Physics, Philosophy, and Polity (Oxford: Clarendon Press), p. 205.
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- ^ Walther Kossel, "Uber Molkulbildung als Frage der Atombau", Ann. Phys., 1916, 49:229-362.
- ^ Pauli admitted in his paper, "On the Connection between the Closing of Electron Groups in Atoms and the Complex Structure of Spectra", published on 21 March 1925 in Zeitschrift für Physik: "We cannot give a more precise reason for this rule." Pais, Abraham (2000), The Genius of Science: A portrait gallery of twentieth-century physicists (New York: Oxford University Press), p. 223.
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Sources
- Gordin, M. D. (2012). "The Textbook Case of a Priority Dispute: D. I. Mendeleev, Lothar Meyer, and the Periodic System". In Biagioli, M.; Riskin, J. (eds.). Nature Engaged. Palgrave Macmillan. pp. 59–82. ISBN 978-1-349-28717-8.
- Mendeleev, D. I. (1870). Естественная система элементов и применение ее к указанию свойств неоткрытых элементов [The natural system of the elements and its application to indication of properties of unknown elements]. pp. 102–176.. Republished from Mendeleev, D. I. (1871). "Естественная система элементовъ и примѣненіе её къ указанію свойствъ неоткрытыхъ элементовъ" [The natural system of the elements and its application to indication of properties of unknown elements]. Journal of the Russian Physico-Chemical Society (in Russian). 3 (2): 25–56. Archived from the original on 17 March 2014.
- Mendeleev, D. I. (1871). Периодическая законность химических элементов [Periodic regularity of the chemical elements]. pp. 102–176.. Republished from Mendelejeff, D. (1871). "Die periodische Gesetzmässigkeit der Elemente" [Periodic regularity of the chemical elements]. Annalen der Chemie und Pharmacie (in German): 133–229.
- Petrov, L. P. (1981). Прогнозирование и размещение инертных элементов в периодической системе [Forecasting and placing of inert elements in the periodic system] (in Russian).
- ISBN 978-0-19-091436-3.
- Thyssen, Pieter; Binnemans, Koen (2015). "Mendeleev and the Rare-Earth Crisis". Philosophy of Chemistry (PDF). Boston Studies in the Philosophy and History of Science. Vol. 306. pp. 155–182. ISBN 978-94-017-9363-6.
Further reading
- Mendeleev, D. I. (1902). Попытка химического понимания мирового эфира [Attempt of chemical understanding of the world ether]. pp. 470–517.. Republished from Mendeleev, D. (1905). Попытка химическаго пониманія мірового эѳира [Attempt of chemical understanding of the world ether] (in Russian). M. P. Frolova's typo-lithography. pp. 5–40.
- Mendeleev, D. I. (1958). Kedrov, K. M. (ed.). Периодический закон [The periodic law] (in Russian). Academy of Sciences of the USSR.
- Trifonov, D. I., ed. (1981). Учение о периодичности: история и современность [Teaching of periodicity: history and modernity] (in Russian). Nauka.
External links
- Development of the periodic table (part of a collection of pages that explores the periodic table and the elements) by the Royal Society of Chemistry
- Dr. Eric Scerri's web page, which contains interviews, lectures and articles on various aspects of the periodic system, including the history of the periodic table.
- The Internet Database of Periodic Tables – a large collection of periodic tables and periodic system formulations.
- History of Mendeleev periodic table of elements as a data visualization at Stack Exchange